Periodic Table of Elements - Trends in The First Ionisation Energy
The periodic table is one of the most important scientific documents known to man. It contains all of the elements currently known to man, along with their element symbol, mass number and atomic number.
These elements are split into groups and periods, which are the columns and rows of the table. Also, the elements are split into three main categories; the first and largest being metals, the second being non-metals and third being the metalloids. These metals are composed 6 subcategories; alkali metals, alkaline earth metals, transition metals, post transition metals, lanthanides and actinides. The non-metals are composed of 3 subcategories; noble gases, halogens and other non-metals.
These categories are substantial, as they group the elements by the bonding they experience, determining other similar properties. There are also many trends within the periodic table, such as boiling and melting points, the atomic radii and the first ionization energy. By looking at the properties of each element, such as its atomic radius, its metallic and non-metallic characteristics, and its atomic mass number, we may determine how these trends appear, where they are and why they are there.
The first ionisation energy is the energy required to remove valence electrons from one mole of gaseous atoms to produce 1 mole of gaseous ions each with a charge of 1+. As the groups increase, the ionisation energy also increases and as the periods increase, the ionisation energy decreases. When talking about ionisation energies, all elements are represented in the gaseous state. These ionisation energies are measured in kJ mol-1, and from about 300 to 2500 is an average to expect them in. All elements have first ionisation energies, though ones that don’t form positive ions naturally require a huge amount of energy to remove one of its electrons.
As shown in the table below, the group 2 elements first ionisation energy decreases as the elements go further down the periodic table.
Element
Proton Number
Symbol
First Ionisation Energy (kJ/mol)
Beryllium
4
Be
900
Magnesium
12
Mg
738
Calcium
20
Ca
590
Strontium
38
Sr
550
Barium
56
Ba
503
First ionisation energy varies in a repetitive manner as you move through the periodic table.
E.g. the pattern from Lithium to neon is identical to the pattern from Sodium to Argon. All of these patterns and variations are a result of the structure of the atoms.
The factors that affect the size of ionisation energy have to do with the attraction between the electron and the nucleus. When the attraction between the electron and nucleus is higher, the ionisation energy is also higher. These factors include the atomic number of the element, the distance of the electron from the nucleus, and whether or not the electron is paired with another or on its own. If the charge on the nucleus is more positive, the electrons are more strongly attracted to it, resulting in a higher attraction.
Also, the further the electron is from the nucleus, the less attracted to the nucleus it is, resulting in lower ionisation energy. As the amount of electrons layered between the nucleus and the outer electron levels increase (also known as screening), the ionisation energy decreases. When there is more than one electron in the same orbital, they experience some repulsion from one another, which slightly counterbalances the attraction, resulting in the electrons being removed rather easily.
Both hydrogen and helium have their outer electrons within the first orbital. As they both have their electrons so close to the nucleus, with no others ‘shielding’ them from it, they have much higher first ionisation energies than that of lithium (hydrogen’s being 1310kJ mol-1 and helium’s being 2370 kJ mol-1). Lithium has its outer electron in the second orbital, and therefore the first orbital is providing the ‘screening’ affect. This causes a reduced pull from the nucleus and lowers ionisation energy (lithium’s being 519kJ mol-1).
Periods two and three have a main trend:
As you can see here, the pattern between the periods are identical, and the only differences are that elements within period two have higher ionisation energies to that of period three.
This general trend applies to all ionisation energies, within a period. As the number of protons in the nucleus increase, the forces of attraction become greater, resulting in greater ionisation energies.
A similar trend applies to the groups on the periodic table. As you move down a group on the periodic table, you notice a decrease in ionisation energies, as shown in the diagram left. As there are less protons within elements higher up in the periodic table, the nuclear charge of the atoms of those further down are much larger. This is due to there being a greater distance to the nucleus, along with more of a ‘screening’ effect.
Finally, we have the trends in ionisation energies of the transition metals. As shown on the right, most of the ionisation energies (excluding zinc, as it’s an outlier), have a ionisation energy of approx. 700kJ mol-1. As the number of protons increase in the transition metals, the 3d electrons also increase, which cancel each other out, resulting in no further attraction from the nucleus. Zinc is the only exception, due to it having one extra proton to increase the attraction. Though there are enough paired electrons in the 4s orbital to cause some repulsion, the one extra proton outweighs them.
Periodic Table of Elements - Trends in The First Ionisation Energy
The periodic table is one of the most important scientific documents known to man. It contains all of the elements currently known to man, along with their element symbol, mass number and atomic number.
These elements are split into groups and periods, which are the columns and rows of the table. Also, the elements are split into three main categories; the first and largest being metals, the second being non-metals and third being the metalloids. These metals are composed 6 subcategories; alkali metals, alkaline earth metals, transition metals, post transition metals, lanthanides and actinides. The non-metals are composed of 3 subcategories; noble gases, halogens and other non-metals.
These categories are substantial, as they group the elements by the bonding they experience, determining other similar properties. There are also many trends within the periodic table, such as boiling and melting points, the atomic radii and the first ionization energy. By looking at the properties of each element, such as its atomic radius, its metallic and non-metallic characteristics, and its atomic mass number, we may determine how these trends appear, where they are and why they are there.
The first ionisation energy is the energy required to remove valence electrons from one mole of gaseous atoms to produce 1 mole of gaseous ions each with a charge of 1+. As the groups increase, the ionisation energy also increases and as the periods increase, the ionisation energy decreases. When talking about ionisation energies, all elements are represented in the gaseous state. These ionisation energies are measured in kJ mol-1, and from about 300 to 2500 is an average to expect them in. All elements have first ionisation energies, though ones that don’t form positive ions naturally require a huge amount of energy to remove one of its electrons.
As shown in the table below, the group 2 elements first ionisation energy decreases as the elements go further down the periodic table.
First ionisation energy varies in a repetitive manner as you move through the periodic table.
E.g. the pattern from Lithium to neon is identical to the pattern from Sodium to Argon. All of these patterns and variations are a result of the structure of the atoms.
The factors that affect the size of ionisation energy have to do with the attraction between the electron and the nucleus. When the attraction between the electron and nucleus is higher, the ionisation energy is also higher. These factors include the atomic number of the element, the distance of the electron from the nucleus, and whether or not the electron is paired with another or on its own. If the charge on the nucleus is more positive, the electrons are more strongly attracted to it, resulting in a higher attraction.
Also, the further the electron is from the nucleus, the less attracted to the nucleus it is, resulting in lower ionisation energy. As the amount of electrons layered between the nucleus and the outer electron levels increase (also known as screening), the ionisation energy decreases. When there is more than one electron in the same orbital, they experience some repulsion from one another, which slightly counterbalances the attraction, resulting in the electrons being removed rather easily.
Both hydrogen and helium have their outer electrons within the first orbital. As they both have their electrons so close to the nucleus, with no others ‘shielding’ them from it, they have much higher first ionisation energies than that of lithium (hydrogen’s being 1310kJ mol-1 and helium’s being 2370 kJ mol-1). Lithium has its outer electron in the second orbital, and therefore the first orbital is providing the ‘screening’ affect. This causes a reduced pull from the nucleus and lowers ionisation energy (lithium’s being 519kJ mol-1).
Periods two and three have a main trend:
This general trend applies to all ionisation energies, within a period. As the number of protons in the nucleus increase, the forces of attraction become greater, resulting in greater ionisation energies.
A similar trend applies to the groups on the periodic table. As you move down a group on the periodic table, you notice a decrease in ionisation energies, as shown in the diagram left. As there are less protons within elements higher up in the periodic table, the nuclear charge of the atoms of those further down are much larger. This is due to there being a greater distance to the nucleus, along with more of a ‘screening’ effect.
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